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C H A P T E R
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T
he United States established its first mint to make silver and gold coins in
Philadelphia in 1792. Some of these old gold and silver coins have become
quite valuable as collector’s items. An 1804 silver dollar recently sold for more
than $4 million. A silver dollar is actually 90% silver and 10% copper. Because
the pure elements gold and silver are too soft to be used alone in coins, other
metals are mixed with them to add strength and durability. These metals
include platinum, copper, zinc, and nickel. Metals make up the majority of the
elements in the periodic table.
START-UPACTIVITY
S A F ET Y P R E C A U T I O N S
What Is a Periodic Table?
PROCEDURE
CONTENTS
SECTION 1
1. Sit in your assigned desk according to the seating chart your
teacher provides.
How Are Elements
Organized?
2. On the blank chart your teacher gives you, jot down information
about yourself—such as name, date of birth, hair color, and
height—in the space that represents where you are seated.
SECTION 2
3. Find out the same information from as many people sitting around
you as possible, and write that information in the corresponding
spaces on the seating chart.
ANALYSIS
1. Looking at the information you gathered, try to identify patterns
that could explain the order of people in the seating chart. If you
cannot yet identify a pattern, collect more information and look
again for a pattern.
4
Tour of the Periodic Table
SECTION 3
Trends in the Periodic Table
SECTION 4
Where Did the Elements
Come From?
2. Test your pattern by gathering information from a person you did
not talk to before.
3. If the new information does not fit in with your pattern, reevaluate
your data to come up with a new hypothesis that explains the
patterns in the seating chart.
Pre-Reading Questions
1
Define element.
2
What is the relationship between the number of protons and
the number of electrons in a neutral atom?
3
As electrons fill orbitals, what patterns do you notice?
www.scilinks.org
Topic: The Periodic Table
SciLinks code: HW4094
115
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S ECTI O N
1
How Are Elements Organized?
KEY TERMS
• periodic law
• valence electron
• group
O BJ ECTIVES
1
Describe the historical development of the periodic table.
2
Describe the organization of the modern periodic table according
to the periodic law.
• period
Patterns in Element Properties
Topic Link
Refer to the chapter “The
Science of Chemistry” for a
definition and discussion
of elements.
Pure elements at room temperature and atmospheric pressure can be
solids, liquids, or gases. Some elements are colorless. Others, like the ones
shown in Figure 1, are colored. Despite the differences between elements,
groups of elements share certain properties. For example, the elements
lithium, sodium, potassium, rubidium, and cesium can combine with chlorine in a 1:1 ratio to form LiCl, NaCl, KCl, RbCl, and CsCl. All of these
compounds are white solids that dissolve in water to form solutions that
conduct electricity.
Similarly, the elements fluorine, chlorine, bromine, and iodine can
combine with sodium in a 1:1 ratio to form NaF, NaCl, NaBr, and NaI.
These compounds are also white solids that can dissolve in water to form
solutions that conduct electricity. These examples show that even though
each element is different, groups of them have much in common.
Figure 1
The elements chlorine,
bromine, and iodine,
pictured from left to right,
look very different from
each other. But each forms
a similar-looking white solid
when it reacts with sodium.
116
Chapter 4
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John Newlands Noticed a Periodic Pattern
Elements vary widely in their properties, but in an orderly way. In 1865,
the English chemist John Newlands arranged the known elements according to their properties and in order of increasing atomic mass. He placed
the elements in a table.
As he studied his arrangement, Newlands noticed that all of the
elements in a given row had similar chemical and physical properties.
Because these properties seemed to repeat every eight elements,
Newlands called this pattern the law of octaves.
This proposed law met with some skepticism when it was first presented, partly because chemists at the time did not know enough about
atoms to be able to suggest a physical basis for any such law.
Dmitri Mendeleev Invented the First Periodic Table
In 1869, the Russian chemist Dmitri Mendeleev used Newlands’s observation and other information to produce the first orderly arrangement, or
periodic table, of all 63 elements known at the time. Mendeleev wrote the
symbol for each element, along with the physical and chemical properties
and the relative atomic mass of the element, on a card. Like Newlands,
Mendeleev arranged the elements in order of increasing atomic mass.
Mendeleev started a new row each time he noticed that the chemical
properties of the elements repeated. He placed elements in the new row
directly below elements of similar chemical properties in the preceding
row. He arrived at the pattern shown in Figure 2.
Two interesting observations can be made about Mendeleev’s table.
First, Mendeleev’s table contains gaps that elements with particular properties should fill. He predicted the properties of the missing elements.
Figure 2
Mendeleev’s table grouped
elements with similar properties into vertical columns.
For example, he placed the
elements highlighted in red
in the table—fluorine,
chlorine, bromine, and
iodine—into the column
that he labeled “VII.”
The Periodic Table
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117
Predicted Versus Actual Properties for Three Elements
Table 1
Properties
Ekaaluminum
Ekaboron
Ekasilicon
(gallium, discovered 1875)
(scandium, discovered 1877)
(germanium, discovered 1886)
Predicted
Observed
Predicted
Observed
Predicted
Observed
6.0 g/cm3
5.96 g/cm3
3.5 g/cm3
3.5 g/cm3
5.5 g/cm3
5.47 g/cm3
low
30ºC
*
*
high
900ºC
Ea2O3
Ga2O3
Eb2O3
Sc2O3
EsO2
GeO2
Solubility of oxide
*
*
*
*
Density of oxide
*
*
*
*
4.7 g/cm3
4.70 g/cm3
Formula of
chloride
*
*
*
*
EsCl4
GeCl4
Color of metal
*
*
*
*
dark gray
grayish white
Density
Melting point
Formula of
oxide
dissolves in acid dissolves in acid
He also gave these elements provisional names, such as “Ekaaluminum”
(the prefix eka- means “one beyond”) for the element that would come
below aluminum. These elements were eventually discovered. As Table 1
illustrates, their properties were close to Mendeleev’s predictions.
Although other chemists, such as Newlands, had created tables of the elements, Mendeleev was the first to use the table to predict the existence of
undiscovered elements. Because Mendeleev’s predictions proved true,
most chemists accepted his periodic table of the elements.
Second, the elements do not always fit neatly in order of atomic mass.
For example, Mendeleev had to switch the order of tellurium, Te, and
iodine, I, to keep similar elements in the same column. At first, he thought
that their atomic masses were wrong. However, careful research by others
showed that they were correct. Mendeleev could not explain why his
order was not always the same.
The Physical Basis of the Periodic Table
About 40 years after Mendeleev published his periodic table, an English
chemist named Henry Moseley found a different physical basis for the
arrangement of elements. When Moseley studied the lines in the
X-ray spectra of 38 different elements, he found that the wavelengths of
the lines in the spectra decreased in a regular manner as atomic mass
increased. With further work, Moseley realized that the spectral lines
correlated to atomic number, not to atomic mass.
When the elements were arranged by increasing atomic number,
the discrepancies in Mendeleev’s table disappeared. Moseley’s work led
to both the modern definition of atomic number, and showed that
atomic number, not atomic mass, is the basis for the organization of the
periodic table.
118
Chapter 4
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1s 1s 1
1s 2
2s 2s 1 2s 2
2p 2p 1 2p 2 2p 3 2p 4 2p 5 2p 6
nd 1–10
3s 3s 1 3s 2
3p 3p 1 3p 2 3p 3 3p 4 3p 5 3p 6
4s 4s 1 4s 2
3d
4p 4p 1 4p 2 4p 3 4p 4 4p 5 4p 6
5s 5s 1 5s 2
4d
5p 5p 1 5p 2 5p 3 5p 4 5p 5 5p 6
6s 6s 1 6s 2
5d
6p 6p 1 6p 2 6p 3 6p 4 6p 5 6p 6
7s 7s 1 7s 2
6d
nf 1–14
5f
6f
Figure 3
The shape of the periodic table is determined by how electrons fill orbitals. Only
the s and p electrons are shown individually because unlike the d and f electrons,
they fill orbitals sequentially.
The Periodic Law
According to Moseley, tellurium, whose atomic number is 52, belongs
before iodine, whose atomic number is 53. Mendeleev had placed these
elements in the same order based on their properties. Today, Mendeleev’s
principle of chemical periodicity is known as the periodic law, which states
that when the elements are arranged according to their atomic numbers,
elements with similar properties appear at regular intervals.
periodic law
Organization of the Periodic Table
valence electron
To understand why elements with similar properties appear at regular
intervals in the periodic table, you need to examine the electron configurations of the elements. As shown in Figure 3, elements in each column of
the table have the same number of electrons in their outer energy level.
These electrons are called valence electrons. It is the valence electrons of
an atom that participate in chemical reactions with other atoms, so elements with the same number of valence electrons tend to react in similar
ways. Because s and p electrons fill sequentially, the number of valence
electrons in s- and p-block elements are predictable. For example, atoms
of elements in the column on the far left have one valence electron.
Atoms of elements in the column on the far right have eight valence
electrons. A vertical column on the periodic table is called a group. A
complete version of the modern periodic table is shown in Figure 4 on the
next two pages.
the law that states that the
repeating physical and chemical
properties of elements change
periodically with their atomic
number
an electron that is found in the
outermost shell of an atom and
that determines the atom’s
chemical properties
group
a vertical column of elements in
the periodic table; elements in a
group share chemical properties
Topic Link
Refer to the chapter “Atoms
and Moles” for a discussion of
electron configuration.
The Periodic Table
Copyright © by Holt, Rinehart and Winston. All rights reserved.
119
Periodic Table of the Elements
1
H
1
Key:
Hydrogen
1.007 94
1s 1
2
Period
3
Group 2
3
4
Average atomic mass
[He]2s 1
[He]2s 2
Electron configuration
12.0107
[He]2s22p2
11
12
Na
Mg
Sodium
22.989 770
Magnesium
24.3050
[Ne]3s 2
Group 3
Group 4
Group 5
Group 6
Group 7
Group 8
Group 9
20
21
22
23
24
25
26
27
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Potassium
39.0983
Calcium
40.078
Scandium
44.955 910
Titanium
47.867
Vanadium
50.9415
Chromium
51.9961
Manganese
54.938 049
Iron
55.845
Cobalt
58.933 200
[Ar]4s 1
[Ar]4s 2
[Ar]3d 14s 2
[Ar]3d 24s 2
[Ar]3d 34s 2
[Ar]3d 54s 1
[Ar]3d 54s 2
[Ar]3d 64s 2
[Ar]3d 74s 2
37
38
39
40
41
42
43
44
45
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Rubidium
85.4678
Strontium
87.62
Yttrium
88.905 85
[Kr]4d 15s 2
Zirconium
91.224
Niobium
92.906 38
Molybdenum
95.94
Technetium
(98)
Ruthenium
101.07
Rhodium
102.905 50
[Kr]4d 25s 2
[Kr]4d 45s1
[Kr]4d 55s 1
[Kr]4d 65s1
[Kr]4d 75s 1
[Kr]4d 85s 1
57
72
73
74
75
76
77
[Kr]5s 2
55
56
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Cesium
132.905 43
Barium
137.327
Lanthanum
138.9055
Hafnium
178.49
Tantalum
180.9479
Tungsten
183.84
Rhenium
186.207
Osmium
190.23
Iridium
192.217
[Xe]6s1
7
Carbon
Be
Beryllium
9.012 182
[Kr]5s1
6
Name
Li
19
5
C
Symbol
Lithium
6.941
[Ne]3s1
4
6
Atomic number
Group 1
[Xe]6s 2
[Xe]5d 16s 2
[Xe]4 f 14 5d 26s 2
[Xe]4f 145d 36s 2
[Xe]4f 145d 46s 2
[Xe]4f 145d 56s 2
[Xe]4f 145d 66s 2
[Xe]4f 145d 76s 2
87
88
89
104
105
106
107
108
109
Fr
Ra
Ac
Rf
Db
Sg
Bh
Hs
Mt
Francium
(223)
Radium
(226)
Actinium
(227)
Rutherfordium
(261)
Dubnium
(262)
Seaborgium
(266)
Bohrium
(264)
Hassium
(277)
Meitnerium
(268)
[Rn]7s 1
[Rn]7s 2
[Rn]6d 17s 2
* The systematic names and symbols
for elements greater than 110 will
be used until the approval of trivial
names by IUPAC.
Topic: Periodic Table
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Keyword: HOLT PERIODIC
Visit the HRW Web site for updates
on the periodic table.
[Rn]5 f 146d 27s 2
[Rn]5 f 146d 37s 2
[Rn]5f 146d 47s 2
[Rn]5 f 146d 57s 2
[Rn]5f 146d 67s 2
[Rn]5 f 146d 77s 2
58
59
60
61
62
Ce
Pr
Nd
Pm
Sm
Cerium
140.116
Praseodymium
140.907 65
Neodymium
144.24
Promethium
(145)
Samarium
150.36
[Xe]4 f 15d 16s 2
[Xe]4 f 36s 2
[Xe]4 f 46s 2
[Xe]4f 56s 2
[Xe]4 f 66s 2
90
91
92
93
94
Th
Pa
U
Np
Pu
Thorium
232.0381
Protactinium
231.035 88
Uranium
238.028 91
Neptunium
(237)
Plutonium
(244)
[Rn]6d 27s 2
[Rn]5f 26d 17s 2
[Rn]5 f 36d 17s 2
[Rn]5 f 46d 17s 2
[Rn]5f 67s 2
Topic: Factors
Affecting
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Chapter 4
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Figure 4
Hydrogen
Semiconductors
(also known as metalloids)
Group 18
2
Metals
Alkali metals
Alkaline-earth metals
Transition metals
Other metals
He
Helium
4.002 602
Nonmetals
Halogens
Noble gases
Other nonmetals
Group 10
Group 11
Group 12
Group 13
Group 14
Group 15
Group 16
Group 17
1s 2
5
6
7
8
9
10
B
C
N
O
F
Ne
Boron
10.811
Carbon
12.0107
Nitrogen
14.0067
Oxygen
15.9994
Fluorine
18.998 4032
Neon
20.1797
[He]2s 22p 1
[He]2s 22p 2
[He]2s 22p 3
[He]2s 22p 4
[He]2s 22p 5
[He]2s 22p 6
13
14
15
16
17
18
Ar
Al
Si
P
S
Cl
Aluminum
26.981 538
Silicon
28.0855
Phosphorus
30.973 761
Sulfur
32.065
Chlorine
35.453
2
[Ne]3s 3p
1
2
[Ne]3s 3p
2
2
[Ne]3s 3p
3
2
[Ne]3s 3p
4
2
[Ne]3s 3p
Argon
39.948
5
[Ne]3s 23p 6
28
29
30
31
32
33
34
35
36
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Nickel
58.6934
Copper
63.546
Zinc
65.409
Gallium
69.723
Germanium
72.64
Arsenic
74.921 60
Selenium
78.96
Bromine
79.904
Krypton
83.798
[Ar]3d 84s 2
[Ar]3d 104s 1
[Ar]3d 104s 2
[Ar]3d 104s 24p 1
[Ar]3d 104s 24p 2
[Ar]3d 104s 24p 3
[Ar]3d 104s 24p 4
[Ar]3d 104s 24p 5
[Ar]3d 104s 24p 6
46
47
48
49
50
51
52
53
54
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Palladium
106.42
Silver
107.8682
Cadmium
112.411
Indium
114.818
Tin
118.710
Antimony
121.760
Tellurium
127.60
Iodine
126.904 47
Xenon
131.293
[Kr]4d 105s 0
[Kr]4d 105s 1
[Kr]4d 105s 2
[Kr]4d 105s 25p 1
[Kr]4d 105s 25p 2
[Kr]4d 105s 25p 3
[Kr]4d 105s 25p 4
[Kr]4d 105s 25p 5
[Kr]4d 105s 25p 6
78
79
80
81
82
83
84
85
86
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
Platinum
195.078
Gold
196.966 55
Mercury
200.59
Thallium
204.3833
Lead
207.2
Bismuth
208.980 38
Polonium
(209)
Astatine
(210)
Radon
(222)
[Xe]4f 145d 96s 1
[Xe]4f 145d 106s 1
[Xe]4f 145d 106s 2
[Xe]4f 145d 106s 26p 1
[Xe]4f 145d 106s 26p 2
[Xe]4f 145d 106s 26p 3
[Xe]4f 145d 106s 26p 4
[Xe]4f 145d 106s 26p 5
[Xe]4f 145d 106s 26p 6
110
111
112
113
114
115
Ds
Uuu*
Uub*
Uut*
Uuq*
Uup*
Darmstadtium
(281)
Unununium
(272)
Ununbium
(285)
Ununtrium
(284)
Ununquadium
(289)
Ununpentium
(288)
[Rn]5f 146d 97s 1
[Rn]5f 146d 107s 1
[Rn]5f 146d 107s 2
[Rn]5f 146d 107s 27p 1
[Rn]5f 146d 107s 27p 2
[Rn]5f 146d 107s 27p 3
A team at Lawrence Berkeley National Laboratories reported the discovery of elements 116 and 118 in June 1999.
The same team retracted the discovery in July 2001. The discovery of elements 113, 114, and 115 has been reported but not confirmed.
63
64
65
66
67
68
69
70
71
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
Europium
151.964
Gadolinium
157.25
Terbium
158.925 34
Dysprosium
162.500
Holmium
164.930 32
Erbium
167.259
Thulium
168.934 21
Ytterbium
173.04
Lutetium
174.967
[Xe]4f 76s 2
[Xe]4f 75d 16s 2
[Xe]4f 96s 2
[Xe]4f 106s 2
[Xe]4f 116s 2
[Xe]4f 126s 2
[Xe]4f 136s 2
[Xe]4f 146s 2
[Xe]4f 145d 16s 2
95
96
97
98
99
100
101
102
103
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Americium
(243)
Curium
(247)
Berkelium
(247)
Californium
(251)
Einsteinium
(252)
Fermium
(257)
Mendelevium
(258)
Nobelium
(259)
Lawrencium
(262)
[Rn]5f 77s 2
[Rn]5f 76d 17s 2
[Rn]5f 97s 2
[Rn]5f 107s 2
[Rn]5f 117s 2
[Rn]5f 127s 2
[Rn]5f 137s 2
[Rn]5f 147s 2
[Rn]5f 146d 17s 2
The atomic masses listed in this table reflect the precision of current measurements. (Values listed in
parentheses are the mass numbers of those radioactive elements’ most stable or most common isotopes.)
The Periodic Table
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121
period
a horizontal row of elements
in the periodic table
1
A horizontal row on the periodic table is called a period. Elements in
the same period have the same number of occupied energy levels. For
example, all elements in Period 2 have atoms whose electrons occupy two
principal energy levels, including the 2s and 2p orbitals. Elements in
Period 5 have outer electrons that fill the 5s, 5d, and 5p orbitals.
This correlation between period number and the number of occupied
energy levels holds for all seven periods. So a periodic table is not needed
to tell to which period an element belongs. All you need to know is the
element’s electron configuration. For example, germanium has the electron configuration [Ar]3d104s24p2. The largest principal quantum number
it has is 4, which means germanium has four occupied energy levels. This
places it in Period 4.
The periodic table provides information about each element, as
shown in the key for Figure 4. This periodic table lists the atomic number,
symbol, name, average atomic mass, and electron configuration in shorthand form for each element.
In addition, some of the categories of elements are designated
through a color code. You may notice that many of the color-coded categories shown in Figure 4 are associated with a certain group or groups.
This shows how categories of elements are grouped by common properties which result from their common number of valence electrons. The
next section discusses the different kinds of elements on the periodic
table and explains how their electron configurations give them their characteristic properties.
Section Review
UNDERSTANDING KEY IDEAS
1. How can one show that elements that have
different appearances have similar chemical
properties?
2. Why was the pattern that Newlands devel-
oped called the law of octaves?
3. What led Mendeleev to predict that some
elements had not yet been discovered?
4. What contribution did Moseley make to the
development of the modern periodic table?
5. State the periodic law.
6. What do elements in the same period have
in common?
7. What do elements in the same group have
in common?
122
CRITICAL THINKING
8. Why can Period 1 contain a maximum of
two elements?
9. In which period and group is the element
whose electron configuration is [Kr]5s1?
10. Write the outer electron configuration for
the Group 2 element in Period 6.
11. What determines the number of elements
found in each period in the periodic table?
12. Are elements with similar chemical
properties more likely to be found in
the same period or in the same group?
Explain your answer.
13. How many valence electrons does
phosphorus have?
14. What would you expect the electron
configuration of element 113 to be?
Chapter 4
Copyright © by Holt, Rinehart and Winston. All rights reserved.
CONSUMER FOCUS
Essential Elements
Table 2
Four elements—hydrogen, oxygen,
carbon, and nitrogen—account for
more than 99% of all atoms in the
human body.
Element
Symbol
Calcium
Ca
Good Health Is Elementary
Hydrogen, oxygen, carbon, and
nitrogen are the major components of the many different
molecules that our bodies need.
Likewise, these elements are the
major elements in the molecules
of the food that we eat.
Another seven elements,
listed in Table 2, are used by our
bodies in substantial quantities,
more than 0.1 g per day. These
elements are known as macronutrients or, more commonly, as
minerals.
Some elements, known as
trace elements or micronutrients,
are necessary for healthy human
Macronutrients
Role in human body chemistry
bones, teeth; essential for blood clotting and
muscle contraction
Phosphorus
P
bones, teeth; component of nucleic acids, including DNA
Potassium
K
present as K+ in all body fluids; essential for nerve action
Sulfur
S
component of many proteins; essential for blood clotting
Chlorine
Cl
present as Cl– in all body fluids; important to
maintaining salt balance
Sodium
Na
present as Na+ in all body fluids; essential for nerve
and muscle action
Magnesium
Mg
in bones and teeth; essential for muscle action
life, but only in very small
amounts. In many cases, humans
need less than 15 nanograms, or
15 × 10–9 g, of a particular trace
element per day to maintain
good health. This means that you
need less than 0.0004 g of such
trace elements during your entire
lifetime!
Questions
1. What do the two macronutrients involved in nerve
action have in common?
2. You may recognize elements
such as arsenic as toxic.
Explain how these elements
can be nutrients even though
they are toxic.
The Periodic Table
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123
S ECTI O N
2
Tour of the Periodic Table
KEY TERMS
• main-group element
O BJ ECTIVES
1
Locate the different families of main-group elements on the periodic
2
Locate metals on the periodic table, describe their characteristic
• alkali metal
• alkaline-earth metal
• halogen
• noble gas
table, describe their characteristic properties, and relate their
properties to their electron configurations.
properties, and relate their properties to their electron configurations.
• transition metal
• lanthanide
• actinide
• alloy
main-group elements
an element in the s-block or
p-block of the periodic table
The Main-Group Elements
Elements in groups 1, 2, and 13–18 are known as the main-group elements.
As shown in Figure 5, main-group elements are in the s- and p-blocks of the
periodic table. The electron configurations of the elements in each main
group are regular and consistent: the elements in each group have the
same number of valence electrons. For example, Group 2 elements have
two valence electrons. The configuration of their valence electrons can be
written as ns2, where n is the period number. Group 16 elements have a
total of six valence electrons in their outermost s and p orbitals. Their
valence electron configuration can be written as ns2np4.
The main-group elements are sometimes called the representative
elements because they have a wide range of properties. At room temperature and atmospheric pressure, many are solids, while others are liquids
or gases. About half of the main-group elements are metals. Many are
extremely reactive, while several are nonreactive. The main-group
elements silicon and oxygen account for four of every five atoms found on
or near Earth’s surface.
Four groups within the main-group elements have special names.
These groups are the alkali metals (Group 1), the alkaline-earth metals
(Group 2), the halogens (Group 17), and the noble gases (Group 18).
Figure 5
Main-group elements
have diverse properties
and uses. They are highlighted in the groups on
the left and right sides
of the periodic table.
124
Main-group
elements
Chapter 4
Copyright © by Holt, Rinehart and Winston. All rights reserved.
Figure 6
The alkali metals make up the first group of
the periodic table. Lithium, pictured here, is an
example of an alkali metal.
The Alkali Metals Make Up Group 1
Elements in Group 1, which is highlighted in Figure 6, are called alkali
metals. Alkali metals are so named because they are metals that react
with water to make alkaline solutions. For example, potassium reacts
vigorously with cold water to form hydrogen gas and the compound
potassium hydroxide, KOH. Because the alkali metals have a single
valence electron, they are very reactive. In losing its one valence electron,
potassium achieves a stable electron configuration.
Alkali metals are usually stored in oil to keep them from reacting with
the oxygen and water in the air. Because of their high reactivity, alkali
metals are never found in nature as pure elements but are found combined with other elements as compounds. For instance, the salt sodium
chloride, NaCl, is abundant in sea water.
Some of the physical properties of the alkali metals are listed in
Table 3. All these elements are so soft that they can be easily cut with a
knife. The freshly cut surface of an alkali metal is shiny, but it dulls quickly
as the metal reacts with oxygen and water in the air. Like other metals, the
alkali metals are good conductors of electricity.
Table 3
alkali metal
one of the elements of Group 1
of the periodic table (lithium,
sodium, potassium, rubidium,
cesium, and francium)
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Topic: Alkali Metals
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Physical Properties of Alkali Metals
Element
Flame test
Hardness
(Mohs’ scale)
Melting
Point (°C)
Boiling
Point (°C)
Density
(g/cm3)
Atomic
radius (pm)
Lithium
red
0.6
180.5
1342
0.53
134
Sodium
yellow
0.4
97.7
883
0.97
154
Potassium
violet
0.5
63.3
759
0.86
196
Rubidium
yellowish violet
0.3
39.3
688
1.53
(216)
Cesium
reddish violet
0.2
28.4
671
1.87
(233)
Refer to Appendix A for more information about the properties of elements, including alkali metals.
The Periodic Table
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125
Figure 7
The alkaline-earth
metals make up the
second group of the
periodic table.
Magnesium, pictured
here, is an example
of an alkaline-earth
metal.
The Alkaline-Earth Metals Make Up Group 2
alkaline-earth metal
one of the elements of Group 2
of the periodic table (beryllium,
magnesium, calcium, strontium,
barium, and radium)
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Topic: Alkaline-Earth Metals
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The Halogens, Group 17, Are Highly Reactive
halogen
one of the elements of Group 17
of the periodic table (fluorine,
chlorine, bromine, iodine, and
astatine); halogens combine with
most metals to form salts
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Topic: Halogens
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126
Group 2 elements, which are highlighted in Figure 7, are called alkalineearth metals. Like the alkali metals, the alkaline-earth metals are highly
reactive, so they are usually found as compounds rather than as pure elements. For example, if the surface of an object made from magnesium is
exposed to the air, the magnesium will react with the oxygen in the air to
form the compound magnesium oxide, MgO, which eventually coats the
surface of the magnesium metal.
The alkaline-earth metals are slightly less reactive than the alkali
metals. The alkaline-earth metals have two valence electrons and must
lose both their valence electrons to get to a stable electron configuration.
It takes more energy to lose two electrons than it takes to lose just the one
electron that the alkali metals must give up to become stable. Although
the alkaline-earth metals are not as reactive, they are harder and have
higher melting points than the alkali metals.
Beryllium is found in emeralds, which are a variety of the mineral
beryl. Perhaps the best-known alkaline-earth metal is calcium, an important mineral nutrient found in the human body. Calcium is essential for
muscle contraction. Bones are made up of calcium phosphate. Calcium
compounds, such as limestone and marble, are common in the Earth’s
crust. Marble is made almost entirely of pure calcium carbonate. Because
marble is hard and durable, it is used in sculptures.
Elements in Group 17 of the periodic table, which are highlighted in
Figure 8 on the next page, are called the halogens. The halogens are the
most reactive group of nonmetal elements because of their electron configuration. Halogens have seven valence electrons—just one short of a
stable configuration. When halogens react, they often gain the one electron needed to have eight valence electrons, a filled outer energy level.
Because the alkali metals have one valence electron, they are ideally
suited to react with the halogens. For example, the alkali metal sodium
easily loses its one valence electron to the halogen chlorine to form the
compound sodium chloride, NaCl, which is table salt. The halogens react
with most metals to produce salts. In fact, the word halogen comes from
Greek and means “salt maker.”
Chapter 4
Copyright © by Holt, Rinehart and Winston. All rights reserved.
Figure 8
The halogens make up Group 17 of the periodic
table. Bromine, one of only two elements that
are liquids at room temperature, is an example
of a halogen.
The halogens have a wide range of physical properties. Fluorine and
chlorine are gases at room temperature, but bromine, depicted in Figure 8,
is a liquid, and iodine and astatine are solids. The halogens are found in sea
water and in compounds found in the rocks of Earth’s crust. Astatine is one
of the rarest of the naturally occurring elements.
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Topic: Noble Gases
SciLinks code: HW4083
The Noble Gases, Group 18, Are Unreactive
Group 18 elements, which are highlighted in Figure 9, are called the
noble gases. The noble gas atoms have a full set of electrons in their
outermost energy level. Except for helium (1s2), noble gases have an
outer-shell configuration of ns2np6. From the low chemical reactivity of
these elements, chemists infer that this full shell of electrons makes these
elements very stable. The low reactivity of noble gases leads to some special uses. Helium, a noble gas, is used to fill blimps because it has a low
density and is not flammable.
The noble gases were once called inert gases because they were
thought to be completely unreactive. But in 1962, chemists were able to
get xenon to react, making the compound XePtF6. In 1979, chemists were
able to form the first xenon-carbon bonds.
noble gas
an unreactive element of Group
18 of the periodic table (helium,
neon, argon, krypton, xenon, or
radon) that has eight electrons in
its outer level (except for helium,
which has two electrons)
Figure 9
The noble gases make up Group 18 of the periodic
table. Helium, whose low density makes it ideal for
use in blimps, is an example of a noble gas.
The Periodic Table
Copyright © by Holt, Rinehart and Winston. All rights reserved.
127
Figure 10
Hydrogen sits
apart from all
other elements in
the periodic table.
Hydrogen is
extremely
flammable and
is used as fuel
for space shuttle
launches.
Hydrogen Is in a Class by Itself
Hydrogen is the most common element in the universe. It is estimated that
about three out of every four atoms in the universe are hydrogen. Because
it consists of just one proton and one electron, hydrogen behaves unlike any
other element. As shown in Figure 10, hydrogen is in a class by itself in the
periodic table.
With its one electron, hydrogen can react with many other elements,
including oxygen. Hydrogen gas and oxygen gas react explosively to form
water. Hydrogen is a component of the organic molecules found in all living things. The main industrial use of hydrogen is in the production of
ammonia, NH3. Large quantities of ammonia are used to make fertilizers.
Most Elements Are Metals
Figure 11 shows that the majority of elements, including many main-group
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Topic: Metals
SciLinks code: HW4079
ones, are metals. But what exactly is a metal? You can often recognize a
metal by its shiny appearance, but some nonmetal elements, plastics,
and minerals are also shiny. For example, a diamond usually has a brilliant
luster. However, diamond is a mineral made entirely of the nonmetal
element carbon.
Conversely, some metals appear black and dull. An example is iron,
which is a very strong and durable metal. Iron is a member of Group 8
and is therefore not a main-group element. Iron belongs to a class of elements called transition metals. However, wherever metals are found on
the periodic table, they tend to share certain properties.
Figure 11
The regions highlighted in
blue indicate the elements
that are metals.
128
Chapter 4
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Metals Share Many Properties
All metals are excellent conductors of electricity. Electrical conductivity
is the one property that distinguishes metals from the nonmetal elements.
Even the least conductive metal conducts electricity 100 000 times better
than the best nonmetallic conductor does.
Metals also exhibit other properties, some of which can also be found
in certain nonmetal elements. For example, metals are excellent conductors
of heat. Some metals, such as manganese and bismuth, are very brittle.
Other metals, such as gold and copper, are ductile and malleable. Ductile
means that the metal can be squeezed out into a wire. Malleable means that
the metal can be hammered or rolled into sheets. Gold, for example, can be
hammered into very thin sheets, called “gold leaf,” and applied to objects
for decoration.
Transition Metals Occupy the Center of the Periodic Table
The transition metals constitute Groups 3 through 12 and are sometimes
called the d-block elements because of their position in the periodic table,
shown in Figure 12. Unlike the main-group elements, the transition
metals in each group do not have identical outer electron configurations.
For example, nickel, Ni, palladium, Pd, and platinum, Pt, are Group 10
metals. However, Ni has the electron configuration [Ar]3d84s2, Pd has the
configuration [Kr]4d 10, and Pt has the configuration [Xe]4f 145d 96s1.
Notice, however, that in each case the sum of the outer d and s electrons
is equal to the group number, 10.
A transition metal may lose different numbers of valence electrons
depending on the element with which it reacts. Generally, the transition
metals are less reactive than the alkali metals and the alkaline-earth
metals are. In fact, some transition metals are so unreactive that they seldom form compounds with other elements. Palladium, platinum, and gold
are among the least reactive of all the elements other than the noble gases.
These three transition metals can be found in nature as pure elements.
Transition metals, like other metals, are good conductors of heat and
electricity. They are also ductile and malleable, as shown in Figure 12.
transition metal
one of the metals that can use
the inner shell before using the
outer shell to bond
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Topic: Transition Metals
SciLinks code: HW4168
Figure 12
Copper, a transition metal, is used in wiring
because it conducts electricity well. Because of
its ductility and malleability, it can be formed
into wires that bend easily.
The Periodic Table
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129
Figure 13
The lanthanides and actinides are placed at the
bottom of the periodic table. Uranium, an actinide,
is used in nuclear reactors. The collection of
uranium-238 kernels is shown here.
Lanthanides and Actinides Fill f-orbitals
lanthanide
a member of the rare-earth
series of elements, whose atomic
numbers range from 58 (cerium)
to 71 (lutetium)
actinide
any of the elements of the
actinide series, which have
atomic numbers from 89
(actinium, Ac) through 103
(lawrencium, Lr)
Part of the last two periods of transition metals are placed toward the
bottom of the periodic table to keep the table conveniently narrow, as
shown in Figure 13. The elements in the first of these rows are called the
lanthanides because their atomic numbers follow the element lanthanum. Likewise, elements in the row below the lanthanides are called
actinides because they follow actinium. As one moves left to right along
these rows, electrons are added to the 4f orbitals in the lanthanides and to
the 5f orbitals in the actinides. For this reason, the lanthanides and
actinides are sometimes called the f-block of the periodic table.
The lanthanides are shiny metals similar in reactivity to the alkalineearth metals. Some lanthanides have practical uses. Compounds of some
lanthanide metals are used to produce color television screens.
The actinides are unique in that their nuclear structures are more
important than their electron configurations. Because the nuclei of
actinides are unstable and spontaneously break apart, all actinides are
radioactive. The best-known actinide is uranium.
Other Properties of Metals
alloy
a solid or liquid mixture of two
or more metals
130
The melting points of metals vary widely. Tungsten has the highest melting point, 4322°C, of any element. In contrast, mercury melts at –39°C, so
it is a liquid at room temperature. This low melting point, along with its
high density, makes mercury useful for barometers.
Metals can be mixed with one or more other elements, usually other
metals, to make an alloy. The mixture of elements in an alloy gives the
alloy properties that are different from the properties of the individual
elements. Often these properties eliminate some disadvantages of the
pure metal. A common alloy is brass, a mixture of copper and zinc, which
is harder than copper and more resistant to corrosion. Brass has a wide
range of uses, from inexpensive jewelry to plumbing hardware. Another
alloy made from copper is sterling silver. A small amount of copper is
mixed with silver to produce sterling silver, which is used for both jewelry
and flatware.
Chapter 4
Copyright © by Holt, Rinehart and Winston. All rights reserved.
Figure 14
Steel is an alloy made of
iron and carbon. When
heated, steel can be worked
into many useful shapes.
Many iron alloys, such as the steel shown in Figure 14, are harder,
stronger, and more resistant to corrosion than pure iron. Steel contains
between 0.2% and 1.5% carbon atoms and often has tiny amounts of
other elements such as manganese and nickel. Stainless steel also incorporates chromium. Because of its hardness and resistance to corrosion,
stainless steel is an ideal alloy for making knives and other tools.
2
Section Review
UNDERSTANDING KEY IDEAS
8. Why are the nuclear structures of the
actinides more important than the electron
configurations of the actinides?
9. What is an alloy?
1. Which group of elements is the most unre-
active? Why?
2. Why do groups among the main-group
elements display similar chemical behavior?
3. What properties do the halogens have in
common?
4. Why is hydrogen set apart by itself?
5. How do the valence electron configurations
of the alkali metals compare with each
other?
6. Why are the alkaline-earth metals less
reactive than the alkali metals?
7. In which groups of the periodic table do
the transition metals belong?
CRITICAL THINKING
10. Noble gases used to be called inert gases.
What discovery changed that term, and
why?
11. If you find an element in nature in its pure
elemental state, what can you infer about
the element’s chemical reactivity?
12. Explain why the transition metals are
sometimes referred to as the d-block
elements.
13. Can an element that conducts heat, is
malleable, and has a high melting point be
classified as a metal? Explain your reasoning.
The Periodic Table
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131
S ECTI O N
3
Trends in the Periodic Table
KEY TERMS
• ionization energy
O BJ ECTIVES
1
Describe periodic trends in ionization energy, and relate them to the
atomic structures of the elements.
2
Describe periodic trends in atomic radius, and relate them to the
atomic structures of the elements.
3
Describe periodic trends in electronegativity, and relate them to the
atomic structures of the elements.
4
Describe periodic trends in ionic size, electron affinity, and melting
and boiling points, and relate them to the atomic structures of the
elements.
• electron shielding
• bond radius
• electronegativity
Periodic Trends
Figure 15
Chemical reactivity with
water increases from top to
bottom for Group 1 elements.
Reactions of lithium, sodium,
and potassium with water
are shown.
The arrangement of the periodic table reveals trends in the properties of
the elements. A trend is a predictable change in a particular direction. For
example, there is a trend in the reactivity of the alkali metals as you move
down Group 1. As Figure 15 illustrates, each of the alkali metals reacts
with water. However, the reactivity of the alkali metals varies. At the top
of Group 1, lithium is the least reactive, sodium is more reactive, and
potassium is still more reactive. In other words, there is a trend toward
greater reactivity as you move down the alkali metals in Group 1.
Understanding a trend among the elements enables you to make predictions about the chemical behavior of the elements. These trends in
properties of the elements in a group or period can be explained in terms
of electron configurations.
Lithium
132
Sodium
Potassium
Chapter 4
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+
–
Electron
lost
Neutral lithium atom
Lithium ion
Li + energy
Li+ + e−
Figure 16
When enough energy is
supplied, a lithium atom
loses an electron and
becomes a positive ion.
The ion is positive because
its number of protons now
exceeds its number of
electrons by one.
Ionization Energy
When atoms have equal numbers of protons and electrons, they are electrically neutral. But when enough energy is added, the attractive force
between the protons and electrons can be overcome. When this happens,
an electron is removed from an atom. The neutral atom then becomes a
positively charged ion.
Figure 16 illustrates the removal of an electron from an atom. The
energy that is supplied to remove an electron is the ionization energy of
the atom. This process can be described as shown below.
A + ionization energy →
neutral atom
A+
ion
+
e−
ionization energy
the energy required to remove
an electron from an atom or ion
electron
Ionization Energy Decreases as You Move Down a Group
Ionization energy tends to decrease down a group, as Figure 17 on the next
page shows. Each element has more occupied energy levels than the one
above it has. Therefore, the outermost electrons are farthest from the
nucleus in elements near the bottom of a group.
Similarly, as you move down a group, each successive element contains more electrons in the energy levels between the nucleus and the
outermost electrons. These inner electrons shield the outermost electrons
from the full attractive force of the nucleus. This electron shielding causes
the outermost electrons to be held less tightly to the nucleus.
Notice in Figure 18 on the next page that the ionization energy of
potassium is less than that of lithium. The outermost electrons of a potassium atom are farther from its nucleus than the outermost electrons of a
lithium atom are from their nucleus. So, the outermost electrons of a
lithium atom are held more tightly to its nucleus. As a result, removing an
electron from a potassium atom takes less energy than removing one
from a lithium atom.
electron shielding
the reduction of the attractive
force between a positively
charged nucleus and its
outermost electrons due to
the cancellation of some of the
positive charge by the negative
charges of the inner electrons
The Periodic Table
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133
Ionization energy
Decreases
Figure 17
Ionization energy generally
decreases down a group and
increases across a period, as
shown in this diagram.
Darker shading indicates
higher ionization energy.
Increases
Ionization Energy Increases as You Move Across a Period
Ionization energy tends to increase as you move from left to right across
a period, as Figure 17 shows. From one element to the next in a period, the
number of protons and the number of electrons increase by one each. The
additional proton increases the nuclear charge. The additional electron is
added to the same outer energy level in each of the elements in the
period. A higher nuclear charge more strongly attracts the outer electrons
in the same energy level, but the electron-shielding effect from inner-level
electrons remains the same. Thus, more energy is required to remove an
electron because the attractive force on them is higher.
Figure 18 shows that the ionization energy of neon is almost four times
greater than that of lithium. A neon atom has 10 protons in its nucleus and
10 electrons filling two energy levels. In contrast, a lithium atom has 3
protons in its nucleus and 3 electrons distributed in the same two energy
levels as those of neon. The attractive force between neon’s 10 protons
and 10 electrons is much greater than that between lithium’s 3 protons
and 3 electrons. As a result, the ionization energy of neon is much higher
than that of lithium.
Ionization Energies of Main-Block Elements
Figure 18
Ionization energies for
hydrogen and for the
main-group elements of the
first four periods are plotted
on this graph.
2400
He
Ne
Ionization energy (kJ/mol)
2000
F
1600
N
H
1200
C
P
Be
800
Mg
Li
400
0
Ca
Na
Al
Si
S
As
Se
15
16
Cl
Kr
Br
Ge
Ga
K
1
B
Ar
O
2
13
14
17
18
Group number
134
Chapter 4
Copyright © by Holt, Rinehart and Winston. All rights reserved.
I2
Br2
Figure 19
In each molecule, half the
distance of the line represents
the bond radius of the atom.
Cl2
Atomic Radius
The exact size of an atom is hard to determine. An atom’s size depends on
the volume occupied by the electrons around the nucleus, and the electrons do not move in well-defined paths. Rather, the volume the electrons
occupy is thought of as an electron cloud, with no clear-cut edge. In addition, the physical and chemical state of an atom can change the size of an
electron cloud.
Figure 19 shows one way to measure the size of an atom. This method
involves calculating the bond radius, the length that is half the distance
between the nuclei of two bonded atoms. The bond radius can change
slightly depending on what atoms are involved.
bond radius
half the distance from center to
center of two like atoms that are
bonded together
Atomic Radius Increases as You Move Down a Group
Atomic radius increases as you move down a group, as Figure 20 shows.
As you proceed from one element down to the next in a group, another
principal energy level is filled. The addition of another level of electrons
increases the size, or atomic radius, of an atom.
Electron shielding also plays a role in determining atomic radius.
Because of electron shielding, the effective nuclear charge acting on the
outer electrons is almost constant as you move down a group, regardless
of the energy level in which the outer electrons are located. As a result,
the outermost electrons are not pulled closer to the nucleus. For example,
the effective nuclear charge acting on the outermost electron in a cesium
atom is about the same as it is in a sodium atom.
Increases
Atomic radius
Figure 20
Atomic radius generally
increases down a group and
decreases across a period, as
shown in this diagram.
Darker shading indicates
higher atomic radius.
Decreases
The Periodic Table
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135
As a member of Period 6, cesium has six occupied energy levels. As a
member of Period 3, sodium has only three occupied energy levels.
Although cesium has more protons and electrons, the effective nuclear
charge acting on the outermost electrons is about the same as it is in
sodium because of electron shielding. Because cesium has more occupied
energy levels than sodium does, cesium has a larger atomic radius than
sodium has. Figure 21 shows that the atomic radius of cesium is about 230
pm, while the atomic radius of sodium is about 150 pm.
Atomic Radius Decreases as You Move Across a Period
Topic Link
Refer to Appendix A for a
chart of relative atomic radii
of the elements.
As you move from left to right across a period, each atom has one more
proton and one more electron than the atom before it has. All additional
electrons go into the same principal energy level—no electrons are being
added to the inner levels. As a result, electron shielding does not play a
role as you move across a period. Therefore, as the nuclear charge
increases across a period, the effective nuclear charge acting on the outer
electrons also increases. This increasing nuclear charge pulls the outermost electrons closer and closer to the nucleus and thus reduces the size
of the atom.
Figure 21 shows how atomic radii decrease as you move across a
period. Notice that the decrease in size is significant as you proceed across
groups going from Group 1 to Group 14. The decrease in size then tends
to level off from Group 14 to Group 18. As the outermost electrons are
pulled closer to the nucleus, they also get closer to one another.
Figure 21
Atomic radii for hydrogen and the main-group
elements in Periods 1
through 6 are plotted
on this graph.
Atomic Radii of Main-Block Elements
250
Cs
Rb
Atomic radius (pm)
200
150
K
Na
Period 6
Period 5
Period 4
Period 3
Li
100
Period 2
50
0
H
1
He
Period 1
2
13
14
15
16
17
18
Group number
136
Chapter 4
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Decreases
Electronegativity
Figure 22
Electronegativity tends to
decrease down a group and
increase across a period, as
shown in this diagram.
Darker shading indicates
higher electronegativity.
Increases
Repulsions between these electrons get stronger. Finally, a point is
reached where the electrons will not come closer to the nucleus because
the electrons would have to be too close to each other. Therefore, the sizes
of the atomic radii level off as you approach the end of each period.
Electronegativity
Atoms often bond to one another to form a compound. These bonds can
involve the sharing of valence electrons. Not all atoms in a compound
share electrons equally. Knowing how strongly each atom attracts bonding electrons can help explain the physical and chemical properties of the
compound.
Linus Pauling, one of America’s most famous chemists, made a scale
of numerical values that reflect how much an atom in a molecule attracts
electrons, called electronegativity values. Chemical bonding that comes
from a sharing of electrons can be thought of as a tug of war. The atom
with the higher electronegativity will pull on the electrons more strongly
than the other atom will.
Fluorine is the element whose atoms most strongly attract shared
electrons in a compound. Pauling arbitrarily gave fluorine an electronegativity value of 4.0. Values for the other elements were calculated in relation to this value.
electronegativity
a measure of the ability of an
atom in a chemical compound
to attract electrons
Electronegativity Decreases as You Move Down a Group
Electronegativity values generally decrease as you move down a group, as
Figure 22 shows. Recall that from one element to the next one in a group,
the principal quantum number increases by one, so another principal
energy level is occupied. The more protons an atom has, the more strongly
it should attract an electron. Therefore, you might expect that electronegativity increases as you move down a group.
However, electron shielding plays a role again. Even though cesium
has many more protons than lithium does, the effective nuclear charge
acting on the outermost electron is almost the same in both atoms. But the
distance between cesium’s sixth principal energy level and its nucleus is
greater than the distance between lithium’s third principal energy level
and its nucleus. This greater distance means that the nucleus of a cesium
atom cannot attract a valence electron as easily as a lithium nucleus can.
Because cesium does not attract an outer electron as strongly as lithium,
it has a smaller electronegativity value.
The Periodic Table
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137
Electronegativity Versus Atomic Number
F
4.0
Period
2
Period
3
Period
4
Period
5
Period
6
3.5
Cl
3.0
Kr
Electronegativity
Xe
2.5
Rn
H
2.0
1.5
1.0
Li
Na
K
Rb
Cs
0.5
0
10
20
30
40
50
60
70
80
Atomic number
Figure 23
This graph shows
electronegativity compared
to atomic number for
Periods 1 through 6.
Electronegativity tends to
increase across a period
because the effective
nuclear charge becomes
greater as protons are
added.
138
Electronegativity Increases as You Move Across a Period
As Figure 23 shows, electronegativity usually increases as you move left to
right across a period. As you proceed across a period, each atom has one
more proton and one more electron—in the same principal energy
level—than the atom before it has. Recall that electron shielding does not
change as you move across a period because no electrons are being added
to the inner levels. Therefore, the effective nuclear charge increases across
a period. As this increases, electrons are attracted much more strongly,
resulting in an increase in electronegativity.
Notice in Figure 23 that the increase in electronegativity across a
period is much more dramatic than the decrease in electronegativity down
a group. For example, if you go across Period 3, the electronegativity more
than triples, increasing from 0.9 for sodium, Na, to 3.2 for chlorine, Cl. In
contrast, if you go down Group 1 the electronegativity decreases only
slightly, dropping from 0.9 for sodium to 0.8 for cesium, Cs.
This difference can be explained if you look at the changes in atomic
structure as you move across a period and down a group. Without the
addition of any electrons to inner energy levels, elements from left to right
in a period experience a significant increase in effective nuclear charge.
As you move down a group, the addition of electrons to inner energy
levels causes the effective nuclear charge to remain about the same. The
electronegativity drops slightly because of the increasing distance
between the nucleus and the outermost energy level.
Chapter 4
Copyright © by Holt, Rinehart and Winston. All rights reserved.
Other Periodic Trends
You may have noticed that effective nuclear charge and electron shielding are often used in explaining the reasons for periodic trends. Effective
nuclear charge and electron shielding also account for two other periodic
trends that are related to the ones already discussed: ionic size and electron affinity. Still other trends are seen by examining how melting point
and boiling point change as you move across a period or down a group.
The trends in melting and boiling points are determined by how electrons
form pairs as d orbitals fill.
Periodic Trends in Ionic Size and Electron Affinity
Recall that atoms form ions by either losing or gaining electrons. Like
atomic size, ionic size has periodic trends. As you proceed down a group,
the outermost electrons in ions are in higher energy levels. Therefore, just
as atomic radius increases as you move down a group, usually the ionic
radius increases as well, as shown in Figure 24a. These trends hold for both
positive and negative ions.
Metals tend to lose one or more electrons and form a positive ion. As
you move across a period, the ionic radii of metal cations tend to decrease
because of the increasing nuclear charge. As you come to the nonmetal
elements in a period, their atoms tend to gain electrons and form negative
ions. Figure 24a shows that as you proceed through the anions on the right
of a period, ionic radii still tend to decrease because of the anions’
increasing nuclear charge.
Neutral atoms can also gain electrons. The energy change that occurs
when a neutral atom gains an electron is called the atom’s electron
affinity. This property of an atom is different from electronegativity, which
is a measure of an atom’s attraction for an electron when the atom is
bonded to another atom. Figure 24b shows that electron affinity tends to
decrease as you move down a group. This trend is due to the increasing
effect of electron shielding. In contrast, electron affinity tends to increase
as you move across a period because of the increasing nuclear charge.
Ionic radii
Decreases
Increases
Cations
Electron affinity
Anions
Cations decrease
Anions decrease
a
Increases
b
Figure 24
Ionic size tends to increase down groups and decrease
across periods. Electron affinity generally decreases
down groups and increases across periods.
The Periodic Table
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139
Periodic Trends in Melting and Boiling Points
The melting and boiling points for the elements in Period 6 are shown in
Figure 25. Notice that instead of a generally increasing or decreasing trend,
melting and boiling points reach two different peaks as d and p orbitals fill.
Cesium, Cs, has low melting and boiling points because it has only one
valence electron to use for bonding. From left to right across the period,
the melting and boiling points at first increase. As the number of electrons
in each element increases, stronger bonds between atoms can form. As a
result, more energy is needed for melting and boiling to occur.
Near the middle of the d-block, the melting and boiling points reach
a peak. This first peak corresponds to the elements whose d orbitals are
almost half filled. The atoms of these elements can form the strongest
bonds, so these elements have the highest melting and boiling points in
this period. For Period 6, the elements with the highest melting and boiling points are tungsten, W, and rhenium, Re.
Melting Points and Boiling Points of Period 6 Elements
Figure 25
As you move across Period
6, the periodic trend for
melting and boiling points
goes through two cycles of
first increasing, reaching a
peak, and then decreasing.
6800
6400
6000
Boiling points
Melting points
5600
5200
4800
Temperature (K)
4400
4000
3600
3200
2800
2400
2000
1600
1200
800
400
0
55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Ti Pb Bi Po At Rn
Atomic number and symbol
140
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As more electrons are added, they begin to form pairs within the d
orbitals. Because of the decrease in unpaired electrons, the bonds that the
atoms can form with each other become weaker. As a result, these elements have lower melting and boiling points. The lowest melting and boiling points are reached at mercury, whose d orbitals are completely filled.
Mercury, Hg, has the second-lowest melting and boiling points in this
period. The noble gas radon, Rn, is the only element in Period 6 with a
lower boiling point than that of mercury.
As you proceed past mercury, the melting and boiling points again
begin to rise as electrons are now added to the p orbital. The melting and
boiling points continue to rise until they peak at the elements whose p
orbitals are almost half filled. Another decrease is seen as electrons pair
up to fill p orbitals. When the noble gas radon, Rn, is reached, the p
orbitals are completely filled. The noble gases are monatomic and have no
bonding forces between atoms. Therefore, their melting and boiling points
are unusually low.
3
Section Review
UNDERSTANDING KEY IDEAS
1. What is ionization energy?
2. Why is measuring the size of an atom difficult?
3. What can you tell about an atom that has
high electronegativity?
4. How does electron shielding affect atomic
size as you move down a group?
5. What periodic trends exist for ionization
energy?
6. Describe one way in which atomic radius is
defined.
7. Explain how the trends in melting and
boiling points differ from the other
periodic trends.
8. Why do both atomic size and ionic size
increase as you move down a group?
9. How is electron affinity different from
electronegativity?
10. What periodic trends exist for
electronegativity?
11. Why is electron shielding not a factor when
you examine a trend across a period?
CRITICAL THINKING
12. Explain why the noble gases have high
ionization energies.
13. What do you think happens to the size of
an atom when the atom loses an electron?
Explain.
14. With the exception of the noble gases, why
is an element with a high ionization energy
likely to have high electron affinity?
15. Explain why atomic radius remains almost
unchanged as you move through Period 2
from Group 14 to Group 18.
16. Helium and hydrogen have almost the same
atomic size, yet the ionization energy of
helium is almost twice that of hydrogen.
Explain why hydrogen has a much higher
ionization energy than any element in
Group 1 does.
17. Why does mercury, Hg, have such a low
melting point? How would you expect
mercury’s melting point to be different if
the d-block contained more groups than
it does?
18. What exceptions are there in the increase
of ionization energies across a period?
The Periodic Table
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141
S ECTI O N
4
Where Did the Elements
Come From?
KEY TERMS
• nuclear reaction
• superheavy element
O BJ ECTIVES
1
Describe how the naturally occurring elements form.
2
Explain how a transmutation changes one element into another.
3
Describe how particle accelerators are used to create synthetic
elements.
Natural Elements
Of all the elements listed in the periodic table, 93 are found in nature.
Three of these elements, technetium, Tc, promethium, Pm, and neptunium, Np, are not found on Earth but have been detected in the spectra
of stars. The nebula shown in Figure 26 is one of the regions in the galaxy
where new stars are formed and where elements are made.
Most of the atoms in living things come from just six elements. These
elements are carbon, hydrogen, oxygen, nitrogen, phosphorus, and sulfur.
Scientists theorize that these elements, along with all 93 natural elements,
were created in the centers of stars billions of years ago, shortly after the
universe formed in a violent explosion.
Figure 26
Three natural elements—
technetium, promethium,
and neptunium—have been
detected only in the spectra
of stars.
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Topic: Origin of Elements
SciLinks code: HW4093
142
Chapter 4
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+
Nuclear
fusion
4 11 H nuclei
4
2 He
nucleus
Figure 27
Nuclear reactions like those in the sun can fuse
four hydrogen nuclei into one helium nucleus,
releasing gamma radiation, .
Hydrogen and Helium Formed After the Big Bang
Much of the evidence about the universe’s origin points toward a single
event: an explosion of unbelievable violence, before which all matter in
the universe could fit on a pinhead. This event is known as the big bang.
Most scientists currently accept this model about the universe’s beginnings. Right after the big bang, temperatures were so high that matter
could not exist; only energy could. As the universe expanded, it cooled
and some of the energy was converted into matter in the form of electrons, protons, and neutrons. As the universe continued to cool, these
particles started to join and formed hydrogen and helium atoms.
Over time, huge clouds of hydrogen accumulated. Gravity pulled
these clouds of hydrogen closer and closer. As the clouds grew more
dense, pressures and temperatures at the centers of the hydrogen clouds
increased, and stars were born. In the centers of stars, nuclear reactions
took place. The simplest nuclear reaction, as shown in Figure 27, involves
fusing hydrogen nuclei to form helium. Even now, these same nuclear
reactions are the source of the energy that we see as the stars’ light and
feel as the sun’s warmth.
nuclear reaction
a reaction that affects the nucleus
of an atom
Other Elements Form by Nuclear Reactions in Stars
The mass of a helium nucleus is less than the total mass of the four hydrogen nuclei that fuse to form it. The mass is not really “lost” in this nuclear
reaction. Rather, the missing mass is converted into energy. Einstein’s
equation E = mc2 describes this mass-energy relationship quantitatively.
The mass that is converted to energy is represented by m in this equation.
The constant c is the speed of light. Einstein’s equation shows that fusion
reactions release very large amounts of energy. The energy released by a
fusion reaction is so great it keeps the centers of the stars at very high
temperatures.
The Periodic Table
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143
Figure 28
Nuclear reactions can form
a beryllium nucleus by fusing
helium nuclei. The beryllium
nucleus can then fuse with
another helium nucleus to form
a carbon nucleus.
+
4
2 He
+
4
2 He
8
4 Be
+
8
4 Be
12
6C
+
+
4
2 He
+
The temperatures in stars get high enough to fuse helium nuclei with
one another. As helium nuclei fuse, elements of still higher atomic numbers
form. Figure 28 illustrates such a process: two helium nuclei fuse to form a
beryllium nucleus, and gamma radiation is released. The beryllium nucleus
can then fuse with another helium nucleus to form a carbon nucleus. Such
repeated fusion reactions can form atoms as massive as iron and nickel.
Very massive stars (stars whose masses are more than 100 times the
mass of our sun) are the source of heavier elements. When such a star has
converted almost all of its core hydrogen and helium into the heavier elements up to iron, the star collapses and then blows apart in an explosion
called a supernova. All of the elements heavier than iron on the periodic
table are formed in this explosion. The star’s contents shoot out into
space, where they can become part of newly forming star systems.
Transmutations
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Topic: Alchemy
SciLinks code: HW4006
In the Middle Ages, many early chemists tried to change, or transmute,
ordinary metals into gold. Although they made many discoveries that
contributed to the development of modern chemistry, their attempts to
transmute metals were doomed from the start. These early chemists did
not realize that a transmutation, whereby one element changes into
another, is a nuclear reaction. It changes the nucleus of an atom and
therefore cannot be achieved by ordinary chemical means.
Transmutations Are a Type of Nuclear Reaction
Although nuclei do not change into different elements in ordinary chemical reactions, transmutations can happen. Early chemists such as John
Dalton had insisted that atoms never change into other elements, so when
scientists first encountered transmutations in the 1910s, their results were
not always believed.
While studying the passage of high-speed alpha particles (helium
nuclei) through water vapor in a cloud chamber, Ernest Rutherford
observed some long, thin particle tracks. These tracks matched the ones
caused by protons in experiments performed earlier by other scientists.
144
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Figure 29
Observe the spot in this
cloud-chamber photo
where an alpha particle
collided with the nucleus
of a nitrogen atom. The left
track was made by an
oxygen atom; the right
track, by a proton.
Proton
(H nucleus)
Alpha particle
Collision
Oxygen ion
Rutherford reasoned correctly that the atomic nuclei in air were disintegrating upon being struck by alpha particles. He believed that the
nuclei in air had disintegrated into the nuclei of hydrogen (protons) plus
the nuclei of some other atom.
Two chemists, an American named W. D. Harkins and an Englishman
named P.M.S. Blackett, studied this strange phenomenon further. Blackett
took photos of 400 000 alpha particle tracks that formed in cloud
chambers. He found that 8 of these tracks forked to form a Y, as shown in
Figure 29. Harkins and Blackett concluded that the Y formed when an
alpha particle collided with a nitrogen atom in air to produce an oxygen
atom and a proton, and that a transmutation had thereby occurred.
Synthetic Elements
The discovery that a transmutation had happened started a flood of
research. Soon after Harkins and Blackett had observed a nitrogen atom
forming oxygen, other transmutation reactions were discovered by bombarding various elements with alpha particles.As a result, chemists have synthesized, or created, more elements than the 93 that occur naturally. These
are synthetic elements. All of the transuranium elements, or those with more
than 92 protons in their nuclei, are synthetic elements. To make them, one
must use special equipment, called particle accelerators, described below.
The Cyclotron Accelerates Charged Particles
Many of the first synthetic elements were made with the help of a
cyclotron, a particle accelerator invented in 1930 by the American scientist E.O. Lawrence. In a cyclotron, charged particles are given one pulse
of energy after another, speeding them to very high energies. The particles
then collide and fuse with atomic nuclei to produce synthetic elements
that have much higher atomic numbers than naturally occurring elements
do. However, there is a limit to the energies that can be reached with a
cyclotron and therefore a limit to the synthetic elements that it can make.
The Periodic Table
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145
The Synchrotron Is Used to Create Superheavy Elements
As a particle reaches a speed of about one-tenth the speed of light, it gains
enough energy such that the relation between energy and mass becomes an
obstacle to any further acceleration. According to the equation E ⫽ mc2,
the increase in the particle’s energy also means an increase in its mass. This
makes the particle accelerate more slowly so that it arrives too late for the
next pulse of energy from the cyclotron, which is needed to make the particle go faster.
The solution was found with the synchrotron, a particle accelerator that
times the pulses to match the acceleration of the particles. A synchrotron
can accelerate only a few types of particles, but those particles it can accelerate reach enormous energies. Synchrotrons are now used in many areas
of basic research, including explorations into the foundations of matter
itself. The Fermi National Accelerator Laboratory in Batavia, IL has a circular accelerator which has a circumference of 4 mi! Subatomic particles
are accelerated through this ring to 99.9999% of the speed of light.
Synthetic Element Trivia
Rutherfordium
Discovered by Russian scientists
at the Joint Institute for Nuclear
Research at Dubna and by
scientists at the University of
California at Berkeley
104
105
106
Meitnerium
Discovered August 29, 1982, by
scientists at the Heavy Ion Research
Laboratory in Darmstadt, West
Germany; named in honor of Lise
Meitner, the Austrian physicist
107
108
109
110
Mendelevium
Synthesized in 1955 by
G. T. Seaborg, A. Ghiorso,
B. Harvey, G. R. Choppin,
and S. G. Thompson at the
University of California,
Berkeley; named in honor
of the inventor of the
periodic system
111
Rf
Db
Sg
Bh
Hs
Mt
Ds
Uuu
Rutherfordium
Dubnium
Seaborgium
Bohrium
Hassium
Meitnerium
Darmstadtium
(unnamed)
93
94
95
96
97
98
99
100
101
102
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Neptunium
Plutonium
Americium
Curium
Berkelium
Californium
Einsteinium
Fermium
Mendelevium
Nobelium
Lawrencium
Curium
Synthesized in 1944
by G. T. Seaborg, R.A.
James, and A. Ghiorso
at the University of
California at Berkeley;
named in honor of
Marie and Pierre Curie
Californium
Synthesized in 1950 by G. T. Seaborg,
S. G. Thompson, A. Ghiorso, and
K. Street, Jr., at the University of
California at Berkeley; named in
honor of the state of California
103
Nobelium
Synthesized in 1958
by A. Ghiorso, G. T. Seaborg,
T. Sikkeland, and J. R. Walton;
named in honor of Alfred Nobel,
discoverer of dynamite and
founder of the Nobel Prize
Figure 30
All of the highlighted elements are synthetic.
Those shown in orange were created by
making moving particles collide with
stationary targets. The elements shown in
blue were created by making nuclei collide.
146
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Once the particles have been accelerated, they are made to collide with
one another. Figure 30 shows some of the superheavy elements created
with such collisions. When a synchrotron is used to create an element, only
a very small number of nuclei actually collide. As a result, only a few nuclei
may be created in these collisions. For example, only three atoms of meitnerium were detected in the first attempt, and these atoms lasted for only
0.0034 s. Obviously, identifying elements that last for such a short time is a
difficult task. Scientists in only a few nations have the resources to carry out
such experiments. The United States, Germany, Russia, and Sweden are the
locations of the largest such research teams.
One of the recent superheavy elements that scientists report is element
114. To create element 114, Russian scientists took plutonium-244, supplied by American scientists, and bombarded it with accelerated
calcium-40 atoms for 40 days. In the end, only a single nucleus was
detected. It lasted for 30 seconds before decaying into element 112.
Most superheavy elements exist for only a tiny fraction of a second.
Thirty seconds is a very long life span for a superheavy element. This long
life span of element 114 points to what scientists have long suspected: that
an “island of stability” would be found beginning with element 114. Based
on how long element 114 lasted, their predictions may have been correct.
However, scientists still must try to confirm that element 114 was in fact
created. The results of a single experiment are never considered valid
unless the experiments are repeated and produce the same results.
4
Section Review
UNDERSTANDING KEY IDEAS
1. How and where did the natural elements
form?
2. What element is the building block for all
other natural elements?
3. What is a synthetic element?
4. What is a transmutation?
5. Why is transmutation classified as a nuclear
reaction?
6. How did Ernest Rutherford deduce that he
had observed a transmutation in his cloud
chamber?
7. How are cyclotrons used to create synthetic
elements?
8. How are superheavy elements created?
superheavy element
an element whose atomic
number is greater than 106
CRITICAL THINKING
9. Why is the following statement not an
example of a transmutation? Zinc reacts
with copper sulfate to produce copper and
zinc sulfate.
10. Elements whose atomic numbers are
greater than 92 are sometimes referred to
as the transuranium elements. Why?
11. Why must an extremely high energy level
be reached before a fusion reaction can
take place?
12. If the synchrotron had not been developed,
how would the periodic table look?
13. What happens to the mass of a particle as
the particle approaches the speed of light?
14. How many different kinds of nuclear
reactions must protons go through to
produce a carbon atom?
The Periodic Table
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147
SCIENCE AND TECHNOLOGY
C A R E E R A P P L I C AT I O N
Superconductors
Superconductivity
Discovered
Materials Scientist
A materials scientist is interested in discovering materials that can last through
harsh conditions, have unusual properties, or perform unique functions. These
materials might include the following: a
lightweight plastic that conducts electricity; extremely light but strong materials
to construct a space platform; a plastic
that can replace iron and aluminum in
building automobile engines; a new
building material that expands and
contracts very little, even in extreme
temperatures; or a strong, flexible, but
extremely tough material that can
replace bone or connective tissue in surgery. Materials engineers develop such
materials and discover ways to mold or
shape these materials into usable forms.
Many materials scientists work in the
aerospace industry and develop new
materials that can lower the mass of
aircraft, rockets, and space vehicles.
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Topic: Superconductors
SciLinks code: HW4170
148
It has long been known that a metal
becomes a better conductor as its
temperature is lowered. In 1911, Heike
Kamerlingh Onnes, a Dutch physicist,
The strong magnetic field
was studying this effect on mercury.
produced by these superconducting electromagnets
When he used liquid helium to cool the
can suspend this 8 cm disk.
metal to about −269°C, an unexpected
thing happened—the mercury lost all resistance and became a
superconductor. Scientists were excited about this new discovery,
but the use of superconductors was severely limited by the huge
expense of cooling them to near absolute zero. Scientists began
research to find a material that would superconduct at temperatures
above −196°C, the boiling point of cheap-to-produce liquid nitrogen.
“High-Temperature” Superconductors
Finally, in 1987 scientists discovered materials that became
superconductors when cooled to only −183°C. These “high-temperature” superconductors were not metals but ceramics; usually
copper oxides combined with elements such as yttrium or barium.
High-temperature superconductors are used in building very
powerful electromagnets that are not limited by resistance or heat.
These magnets can be used to build powerful particle accelerators
and high-efficiency electric motors and generators. Engineers are
working to build a system that will use superconducting electromagnets to levitate a passenger train above its guide rail so that the
train can move with little friction and thus save fuel.
Questions
1. How does temperature normally affect electrical conductivity
in metals?
2. What happened unexpectedly when mercury was cooled to
near absolute zero?
3. How might consumers benefit from the use of superconducting
materials?
Chapter 4
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CHAPTER HIGHLIGHTS
KEY TERMS
periodic law
valence electron
group
period
main-group element
alkali metal
alkaline-earth metal
halogen
noble gas
transition metal
lanthanide
actinide
alloy
ionization energy
electron shielding
bond radius
electronegativity
nuclear reaction
superheavy element
4
KEY I DEAS
SECTION ONE How Are Elements Organized?
• John Newlands, Dmitri Mendeleev, and Henry Moseley contributed to the
development of the periodic table.
• The periodic law states that the properties of elements are periodic functions
of the elements’ atomic numbers.
• In the periodic table, elements are ordered by increasing atomic number. Rows
are called periods. Columns are called groups.
• Elements in the same period have the same number of occupied energy levels.
Elements in the same group have the same number of valence electrons.
SECTION TWO Tour of the Periodic Table
• The main-group elements are Group 1 (alkali metals), Group 2 (alkaline-earth
metals), Groups 13–16, Group 17 (halogens), and Group 18 (noble gases).
• Hydrogen is in a class by itself.
• Most elements are metals, which conduct electricity. Metals are also ductile
and malleable.
• Transition metals, including the lanthanides and actinides, occupy the center
of the periodic table.
SECTION THREE Trends in the Periodic Table
• Periodic trends are related to the atomic structure of the elements.
• Ionization energy, electronegativity, and electron affinity generally increase as
you move across a period and decrease as you move down a group.
• Atomic radius and ionic size generally decrease as you move across a period
and increase as you move down a group.
• Melting points and boiling points pass through two cycles of increasing,
peaking, and then decreasing as you move across a period.
SECTION FOUR Where Did the Elements Come From?
• The 93 natural elements were formed in the interiors of stars. Synthetic
elements (elements whose atomic numbers are greater than 93) are made
using particle accelerators.
• A transmutation is a nuclear reaction in which one nucleus is changed into
another nucleus.
The Periodic Table
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149
4
CHAPTER REVIEW
USING KEY TERMS
UNDERSTANDING KEY IDEAS
1. What group of elements do Ca, Be, and Mg
belong to?
14. How was Moseley’s arrangement of the
2. What group of elements easily gains one
valence electron?
3. What category do most of the elements of
the periodic table fall under?
when an atom gains an electron?
16. Why was Mendeleev’s periodic table
17. What determines the horizontal arrange-
6. Give an example of a nuclear reaction.
Describe the process by which it takes place.
7. What are elements in the first group of the
ment of the periodic table?
18. Why is barium, Ba, placed in Group 2 and
in Period 6?
Tour of the Periodic Table
periodic table called?
8. What atomic property affects periodic
trends down a group in the periodic table?
9. What two atomic properties have an
increasing trend as you move across a
period?
WRITING
SKILLS
your own words how synthetic
elements are created. Discuss
what modification has to be made to the
equipment in order to synthesize superheavy elements.
11. Which group of elements has very high
ionization energies and very low electron
affinities?
12. How many valence electrons does a fluorine
13. Give an example of an alloy.
15. What did the gaps on Mendeleev’s periodic
accepted by most chemists?
5. What are elements 90–103 called?
atom have?
elements in the periodic table different
from Mendeleev’s?
table represent?
4. What is the term for the energy released
10. Write a paragraph describing in
How Are Elements Organized?
19. Why is hydrogen in a class by itself?
20. All halogens are highly reactive. What
causes these elements to have similar
chemical behavior?
21. What property do the noble gases share?
How do the electron configurations of the
noble gases give them this shared property?
22. How do the electron configurations of the
transition metals differ from those of the
metals in Groups 1 and 2?
23. Why is carbon, a nonmetal element, added
to iron to make nails?
24. If an element breaks when it is struck with a
hammer, could it be a metal? Explain.
25. Why are the lanthanides and actinides
placed at the bottom of the periodic table?
26. Explain why the main-group elements are
also known as representative elements.
150
Chapter 4
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Trends in the Periodic Table
Where Did the Elements Come From?
27. What periodic trends exist for ionization
34. How does nuclear fusion generate energy?
energy? How does this trend relate to
different energy levels?
28. Why don’t chemists define atomic radius as
the radius of the electron cloud that surrounds a nucleus?
29. How does the periodic trend of atomic
radius relate to the addition of electrons?
30. What happens to electron affinity as you
move across a period beginning with Group
1? Why do these values change as they do?
31. Identify which trend diagram below
describes atomic radius.
Increases
36. Why are technetium, promethium, and
neptunium considered natural elements
even though they are not found on Earth?
37. Why must a synchrotron be used to create
a superheavy element?
38. What role did supernovae play in creating
the natural elements?
39. What two elements make up most of the
matter in a star?
MIXED REVIEW
identify the period and group in which
each of the following elements is located.
1
a. [Rn]7s
2
b. [Ar]4s
2
6
c. [Ne]3s 3p
Decreases
41. Which of the following ions has the electron
configuration of a noble gas: Ca+ or Cl−?
(Hint: Write the electron configuration for
each ion.)
Decreases
b.
when a transmutation takes place?
40. Without looking at the periodic table,
Increases
a.
35. What happens in the nucleus of an atom
42. When 578 kJ/mol of energy is supplied, Al
Decreases
Increases
c.
loses one valence electron. Write the
electron configuration of the ion that forms.
32. What periodic trends exist for electronega-
tivity? Explain the factors involved.
33. Why are the melting and boiling points of
mercury almost the lowest of the elements
in its period?
43. Name three periodic trends you encounter
in your life.
44. How do the electron configurations of the
lanthanide and actinide elements differ from
the electron configurations of the other
transition metals?
45. Use the periodic table to describe the chem-
ical properties of the following elements:
a. iodine, I
b. krypton, Kr
c. rubidium, Rb
The Periodic Table
Copyright © by Holt, Rinehart and Winston. All rights reserved.
151
46. The electron configuration of argon differs
from those of chlorine and potassium by
one electron each. Compare the reactivity of
these three elements, and relate them to
their electron configurations.
in the periodic table. Does strontium share
more properties with yttrium or barium?
Explain your answer.
54. Examine the following diagram.
47. What trends were first used to classify the
elements? What trends were discovered
after the elements were classified in the
periodic table?
48. Among the main-group elements, what is
the relationship between group number and
the number of valence electrons among
group members?
Explain why the structure shown on the
right was drawn to have a smaller radius
than the structure on the left.
CRITICAL THINKING
49. Consider two main-group elements, A and
B. Element A has an ionization energy of
419 kJ/mol. Element B has an ionization
energy of 1000 kJ/mol. Which element is
more likely to form a cation?
50. Argon differs from both chlorine and potas-
sium by one proton each. Compare the electron configurations of these three elements
to explain the reactivity of these elements.
51. While at an amusement park, you inhale
helium from a balloon to make your voice
higher pitched. A friend says that helium
reacts with and tightens the vocal cords to
make your voice have the higher pitch.
Could he be correct? Why or why not?
52. In his periodic table, Mendeleev placed Be,
Mg, Zn, and Cd in one group and Ca, Sr, Ba,
and Pb in another group. Examine the electron configurations of these elements, and
explain why Mendeleev grouped the elements this way.
53. The atomic number of yttrium, which fol-
lows strontium in the periodic table, exceeds
the atomic number of strontium by one.
Barium is 18 atomic numbers after strontium but it falls directly beneath strontium
152
ALTERNATIVE ASSESSMENT
55. Select an alloy. You can choose one men-
tioned in this book or find another one by
checking the library or the Internet. Obtain
information on how the alloy is made.
Obtain information on how the alloy is
used for practical purposes.
56. Construct a model of a synchrotron. Check
the library and Internet for information
about synchrotrons. You may want to
contact a synchrotron facility directly to
find out what is currently being done in
the field of synthetic elements.
57. In many labeled foods, the mineral content
is stated in terms of the mass of the element,
in a stated quantity of food. Examine the
product labels of the foods you eat.
Determine which elements are represented
in your food and what function each element serves in the body. Make a poster of
foods that are good sources of minerals that
you need.
CONCEPT MAPPING
58. Use the following terms to create a concept
map: atomic number, atoms, electrons, periodic table, and protons.
Chapter 4
Copyright © by Holt, Rinehart and Winston. All rights reserved.
FOCUS ON GRAPHING
Study the graph below, and answer the questions that follow.
For help in interpreting graphs, see Appendix B, “Study Skills for Chemistry.”
59. What relationship is represented in the
Atomic Radii of Main-Block Elements
graph shown?
250
60. What do the numbers on the y-axis
Cs
Rb
represent?
200
K
the element with the greatest atomic
radius?
62. Why is the axis representing group
number drawn the way it is in going
from Group 2 to Group 13?
63. Which period shows the greatest change
Atomic radius (pm)
61. In every Period, which Group contains
150
Na
Period 6
Period 5
Period 4
Period 3
Li
100
Period 2
50
H
He
Period 1
in atomic radius?
64. Notice that the points plotted for the
elements in Periods 5 and 6 of Group 2
overlap. What does this overlap indicate?
0
1
2
13
14
15
16
17
18
Group number
TECHNOLOGY AND LEARNING
65. Graphing Calculator
Graphing Atomic Radius Vs. Atomic Number
The graphing calculator can run a program
that graphs data such as atomic radius versus atomic number. Graphing the data
within the different periods will allow you
to discover trends.
Go to Appendix C. If you are using a TI-83
Plus, you can download the program and
data sets and run the application as directed.
Press the APPS key on your calculator, then
choose the application CHEMAPPS. Press 8,
then highlight ALL on the screen, press 1,
then highlight LOAD and press 2 to load the
data into your calculator. Quit the application, and then run the program RADIUS. For
L1, press 2nd and LIST, and choose ATNUM.
For L2, press 2nd and LIST and choose
ATRAD.
If you are using another calculator, your
teacher will provide you with keystrokes and
data sets to use.
a. Would you expect any atomic number to
have an atomic radius of 20 pm? Explain.
b. A relationship is considered a function if it
can pass a vertical line test. That is, if a vertical line can be drawn anywhere on the graph
and only pass through one point, the relationship is a function. Does this set of data represent a function? Explain.
c. How would you describe the graphical relationship between the atomic numbers and
atomic radii?
The Periodic Table
Copyright © by Holt, Rinehart and Winston. All rights reserved.
153
4
STANDARDIZED TEST PREP
UNDERSTANDING CONCEPTS
Directions (1–4): For each question, write on a
separate sheet of paper the letter of the correct
answer.
1
2
3
Which of the following elements is formed
in stars?
A. curium
C. gold
B. einsteinium
D. mendelevium
Why are the Group 17 elements, the
halogens, the most reactive of the nonmetal
elements?
F. They have the largest atomic radii.
G. They have the highest ionization energies.
H. They are the farthest right on the periodic
table.
I. They require only one electron to fill their
outer energy level.
Which of the following is a property of
noble gases as a result of their stable electron
configuration?
A. large atomic radii
B. high electron affinities
C. high ionization energies
D. a tendency to form both cations
and anions
4 Which of these is a transition element?
F. Ba
H. Fe
G. C
I. Xe
Directions (5–7): For each question, write a short
response.
5
How did the discovery of the elements that
filled the gaps in Mendeleev’s periodic table
increase confidence in the periodic table?
6
Why is iodine placed after tellurium on the
periodic table if the atomic mass of tellurium
is less than that of iodine?
154
7
What is the outermost occupied energy level
in atoms of the elements in Period 4?
READING SKILLS
Directions (8–10): Read the passage below. Then
answer the questions.
The atomic number of beryllium is one less
than that of boron, which follows it on the periodic table. Strontium, which is directly below
beryllium in period 5 of the periodic table has
34 more protons and 34 more electrons than
beryllium. However, the properties of beryllium
resemble the much larger strontium more than
those of similar-sized boron.
8
The properties of beryllium are more similar
to those of strontium than those of boron
because
A. A strontium atom is larger than a boron
atom.
B. Strontium and beryllium are both reactive
nonmetals.
C. A strontium atom has more electrons
than a boron atom.
D. Strontium has the same number of
valence electrons as beryllium.
9
Beryllium and strontium are both located in
the second column of the periodic table. To
which of these classifications do they belong?
F. alkali metals
G. alkaline earth metals
H. rare earth metals
I. transition metals
0
Why is it easier to determine to which column of the periodic table an element belongs
than to determine to which row it belongs,
based on observations of its properties?
Chapter 4
Copyright © by Holt, Rinehart and Winston. All rights reserved.
INTERPRETING GRAPHICS
Directions (11–13): For each question below, record the correct answer on a
separate sheet of paper.
Use the diagram below to answer question 11.
+
+
4
2 He
q
4
2 He
␥
+
8
4 Be
What process is represented by this illustration?
A. chemical reaction
B. ionization
C. nuclear fission
D. nuclear fusion
The graph below shows the ionization energies (kilojoules per mole) of mainblock elements. Use it to answer questions 12 and 13.
Ionization Energies of Main-Block Elements
2400
He
Ne
Ionization energy (kJ/mol)
2000
F
1600
N
H
1200
C
Be
800
Mg
Li
400
0
Ca
Na
Si
Al
Ge
S
As
Se
15
16
Cl
Kr
Br
Ga
K
1
B
P
Ar
O
2
13
14
17
18
Group number
w
e
Which of these elements requires the most energy to remove an electron?
F. argon
H. nitrogen
G. fluorine
I. oxygen
Explain the trend in ionization energy within a group on the periodic
table.
Test
Before looking at the
answer choices for
a question, try to
answer the question
yourself.
Standardized Test Prep
Copyright © by Holt, Rinehart and Winston. All rights reserved.
155